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6-3-1 Compare and Contrast a Chemical Formula for a Molecular Compound with One for an Ionic Compound

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6-3-1 Compare and Contrast a Chemical Formula for a Molecular Compound with One for an Ionic Compound Chemistry Ch 6 sect 3 «F_Name» «L_Name» Period «Per» «num» 6-3-1 Compare and contrast a chemical formula for a molecular compound with one for an ionic compound. Bond: Attraction between 2 or more atoms or ions. Bonding occurs because it lowers the energy of the system. Chemical Bonding is also referred to as valence bond Bonding involves electrons in valence shell (outermost electron shell; has highest principal quantum number (n) Electrons on atom that are not valence electrons are called core electrons Two broad classifications Ionic (attraction between cations & anions; Metal- NonMetal) Covalent (electron sharing; NonMetal - NonMetal) Caution! The idea of a "pure ionic bond" is an oversimplification, even if you consider very strongly ionic species such as NaCl or NaF Ionic bond is the attraction between cations and anions with a lowest energy arrangement (like covalent bonds, the lowest potential is sought). The pattern is repeated throughout the crystal or Lattice Also see lattices on page 177 of text. 1 «F_Name» «L_Name» Room # An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. A formula unit is the simplest collection of atoms from which an ionic compound’s formula can be established. It represents the ratio of atoms in an ionic bond. For example: sodium chloride, the ions (Na+ and Cl - ions) combine in a one-to-one ratio the formula unit is written simply as NaCl (the ones are understood). In Calcium Fluoride (Ca2+ and F – ions) the ions combine in a 1 to 2 ratio, the formula unit is CaF2 . (see fig 6 – 10 page 170 of text). Ions form to reach an octet The Octet Rule: Cations: Group 1 elements lose 1 electron and form cations with +1 charge X + example Na + Group 2 elements lose 2 electrons and form cations with +2 charge 2+ 2+ X example Mg Some group 13 elements lose 3 electrons and form cations with +3 charge X3+ example Al 3+ The formation of Group 1, 2 , and selected 13 cations results in Noble gas electron configuration, when their valence electrons are given away, they acquire the outer shell configuration of the Nobel gas one level lower. Mg 2+ has the electron configuration 1s2 2s2 2p6 just like Neon. All of group 1 and group 2 ions will follow this example. Some group 13 elements will loose d orbital electrons (Ex Sn), thus may not achieve NG Configuration. The Octet Rule: Anions: Group15 (-3), 16 (-2) and 17 (-1) gain electrons to reach their NG configuration Lewis formula for ionic compounds metal-nonmetal Consider NaCl; it consists of an array of Na+ and Cl- ions. Na+ is the Lewis symbol for sodium ions. There are no dots on the symbol because Na has given away it’s valence electrons Note! Cl – has 8 valence electrons so when it is drawn it must have 8 dots around it. It also has a negative charge, as a result anion Lewis dot looks a little different. Square brackets are place on both sides of the ion and the charge is expressed as a superscript: 2 Chemistry Ch 6 sect 3 «F_Name» «L_Name» Period «Per» «num» NaCl represented as + – Na Cl 6-3-2 Discuss the arrangements of ions in crystals. In an ionic crystal, ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice. A crystal lattice is a regularly repeating pattern throughout the compound. The strengths of attraction between ions vary with the sizes and charges of the ions and the numbers of ions of different charges. For example, in calcium fluoride, there are two anions for each cation. 6-3-3 Define lattice energy and explain its significance. Lattice energy is the energy released when one mole of an ionic crystalline compound is formed from gaseous ions. Energy is released when the crystals are formed. The crystal is lower in potential energy and therefore is more stable which makes this form most desirable for the atoms. 6-3-4 List and compare the distinctive properties of ionic and molecular compounds. Compare and contrast Molecular (covalent) and Ionic Bonding The forces of attraction between molecules, known as intermolecular forces are weaker than the forces of ionic bonding. Molecular compounds melt and boil at lower temperatures than ionic compounds because of these weaker forces. Molecular compounds vaporize at room temperature more easily than ionic compounds. Ionic compounds are hard, but they are brittle. Many ionic compounds are soluble in water. When they are dissolved in water, they are electrical conductors. Many purely molecular compounds are not conductors. Caution: many compounds in reality have some ionic character and some molecular character. Some even have both (soap for example in an organic salt, most medications are also organic salts). 3 «F_Name» «L_Name» Room # 6-3-5 Write the Lewis structure for a polyatomic ion given the identity of the atoms combined and other appropriate information. Polyatomic Ions: A charged group of covalently bonded atoms. Like monatomic ions (single atom ion) the charge of a polyatomic ion results from an excess of electrons (negative charge) or a shortage of electrons (positive charge). To find the Lewis structure for a polyatomic ion, follow the 6 steps outlined for covalent bonding, see Sample Problem 6-4 on page 174, with the following exception. If the ion is negatively charged, add to the total number of valence electrons a number corresponding to the ion’s negative charge; # of Val e–s + the absolute value of the negative charge to equal the number of dots. If the ion is positively charged, subtract from the total number of valence electrons a number corresponding to the ion’s positive charge; # of Val e–s – the value of the positive ion’s charge to equal the number of dots necessary. Then square bracket the ion and indicate its charge. Examples Homework: p. 180. Do #’s 1-4; p. 195. Assign #s 15-19, 36, and 42. And WS 4 Chemistry Ch 6 sect 3 «F_Name» «L_Name» Period «Per» «num» 1. Two possible examples are sodium chloride, NaCl, and magnesium chloride, MgCl2. 3. Ionic compounds consist of positive and negative ions bound together by electrical attraction. Molecular compounds are groups of atoms held together by covalent bonding, or the sharing of electrons.. 4. Compound B is probably a molecular substance; compound A is probably ionic. The intermolecular attractions that hold molecules together are weaker than ionic attraction, resulting in lower melting and boiling points in molecular substances. 15. a. An ionic compound is composed of cations and anions such that the total positive and negative charges are equal. b. Most ionic compounds occur naturally as crystalline solids. 16. a. the simplest collection of atoms from which an ionic compound's formula can be established b. one calcium ion, Ca2+, and two fluorine ions, F- 17. a. the energy released when one mole of an ionic compound is formed from gaseous ions b. The greater the lattice energy, the stronger the ionic bonding. 18. a. Ionic compounds have higher melting and boiling points than molecular compounds do, and they do not vaporize at room temperature. 5 «F_Name» «L_Name» Room # b. The differences in the proper-ties of ionic and molecular compounds are generally a result of differences in how strongly the compound's basic units are held together. c. hardness, brittleness, electrical conductivity in the molten state 19. a. a charged group of covalently bonded atoms b. Some common examples of polyatomic anions include the nitrate - + 2- ion, NO3 , the ammonium ion, NH4 , the sulfate ion, SO4 , and the 3- phosphate ion, PO4 . c. Polyatomic ions combine with ions of opposite charge to form ionic compounds. 36. Bonding is stronger between the ions in sodium chloride because its lattice energy is greater (more negative). Greater lattice energy indicates stronger ionic bonding. 42. Draw Lewis structures for each of the following polyatomic ions. Show resonance structures, if they exist. a. OH− − b. H3C2O2 − c. BrO3 6 .
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